# Molar Mass of Butane Lab (PDF)

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Molar Mass of Butane Lab
Purpose: To determine the molar mass of butane

Pre-lab Discussion: In this experiment you will collect a volume of butane gas and use the ideal gas law
equation (PV = nRT) to determine the molecular mass of the gas. In this equation P represents pressure, V is
volume, n is the number of moles of the gas, R is the ideal gas constant, and T is the Kelvin temperature. The
number of moles (n) can be rewritten as being equal to the mass of the gas (m) divided by the molar mass of the
gas (M). A little algebra enables us to rewrite the equation as M = mRT/PV. There are many values of the gas
constant R since there are many allowable units for pressure and volume, but one value is 62.4 x 103 torr–ml /
mole-K.

The pressure of the gas will be adjusted to atmospheric pressure and the mass of the butane will be measured.
Our source of butane will be a cigarette lighter and we will collect the gas by water displacement. This is
possible because butane is not very soluble in water. Dalton’s Law of Partial Pressures will enable us to
determine the pressure of the butane gas alone in the cylinder.

Procedure

1. Weigh the lighter to the nearest 0.01 grams or better. Attach a plastic tube to the gas opening on the top of the
lighter.

2. Submerge a 100. ml graduated cylinder in water so that the cylinder fills completely with water. Invert the
cylinder. Make sure there are no air bubbles remaining in the graduated cylinder.

3. Take the plastic tube and insert it in the graduated cylinder under the water. Carefully release the butane from
the lighter and collect it in the cylinder. Release enough butane to fill the tube to within 5 ml of its graduated
capacity. (95 ml for a 100 ml cylinder). Remove the plastic tube from the graduated cylinder and the lighter.

4. Allow the butane in the cylinder to reach room temperature (about five minutes). Then adjust the level of the
water inside and outside the tube to be the same. With the pressure inside the same as the pressure outside,
record the volume to the nearest ml. This makes the pressure of the combined water vapor and butane gas equal
to the pressure of the atmosphere.

5. Measure the mass of the lighter again.

6. Record the air temperature in Kelvin.

7. Record the water temperature and barometric pressure.

8. Use the derivation of the ideal gas law equation (M = mRT/PV) to find the molar mass.

9. Repeat the entire procedure 2 more times for a total of 3 trials. Find the average and the % error.

rr g:\files\courses\101-2lab\butanelabrev.doc                                         07/05/04
For the quiz:

1. Know how to do the calculations and understand why you did each part of the procedure.

2. Think about how various errors would affect your final result. For example, if the volume of the gas was
high, when substituted into the equation to determine the molar mass, your final answer would be low. Be able
to trace other errors in the same way.

Data Sheet
Trial 1            Trial 2          Trial 3
Mass of lighter assembly and
contents before collection
Mass of the lighter assembly and
contents after collection
Mass of the butane gas evolved

Volume of gas produced at room
temperature and atmospheric pressure
Barometric pressure

Water temperature

Vapor pressure of water at water
temperature (from table)
Partial pressure of dry butane (show
calculation)
Ptot = Pbutane + Pwater vapor

Room temperature

Molar mass of the butane (show
calculation)
M = mRT/PV

True value for the molar mass _______________________

Average molar mass for the three trials _______________

Experimental error for the average (show calculation) __________________

rr g:\files\courses\101-2lab\butanelabrev.doc                                     07/05/04

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